The activation energy for a reaction is $9.0 \mathrm{~K} \mathrm{cal} \mathrm{mol}^{-1}$. The increase in the rate constant when its temperature is increased from 298 K to 308 K is

$A(g) \longrightarrow B(g)$ is a first order reaction. The initial concentration of $A$ is $0.2 \mathrm{~mol} \mathrm{~L}^{-1}$. After 10 minutes the concentration of $B$ is found to be $0.18 \mathrm{~mol} \mathrm{~L}^{-1}$. The rate constant (in $\min ^{-1}$ ) for the reaction is

Which is correct about zero order reaction?

Arrhenius equation may not be represented as

Consider the reaction,

$$2 \mathrm{~N}_2 \mathrm{O}_5(g) \longrightarrow 4 \mathrm{NO}_2(g)+\mathrm{O}_2(g)$$

The rate law for this reaction is rate $$=k\left[\mathrm{~N}_2 \mathrm{O}_5\right]$$.

Which of the following statements is true regarding the above reaction?